Download C5H12 + 8 O2 → 5 CO2 + 6 H2O

Chemical Reactions
Week 12: Lectures 34 – 36
Lecture 34: W 11/9
Lecture 35: F 11/11
Lecture 36: M 11/14
Reading:
BLB Ch 3.1 – 3.2; 3.6 – 3.7; 4.2 – 4.4
Homework:
BLB 3: 1, 64; 4: 24, 39; Supp Rxns: 1 – 11
Reminder:
No Angel Quiz on Thur 11/10
ALEKS Objective 12 due on Tues 11/15
Jensen Office Hour: 501 Chemistry Building
Tuesdays and Thursdays 10:30 – 11:30 am
Late drop deadline: Friday 11/11 @11:59 pm
Final Exam: Monday Dec 12 2:30 – 4:20 pm
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4,5',)6$7&8($98&($:+;,0$<=$>-*($?@A,#'0,$$
Law of conservation of mass: total mass does not
change during a chemical reaction
• Mass of reactants MUST equal mass of products
# of atoms of each element on reactant side
= # of atoms of each element on product side
Example: Complete combustion of pentane
C5H12 + 8 O2
!
5 CO2 + 6 H2O
If mass of reactants is 100 g, then the mass of
products must be ______ g.
Reactants: ____C, ____H, ____O atoms
Products: ____C, ____H, ____O atoms
___ mol C5H12 react completely with ____ mol O2,
produce ____ mol CO2 and ____ mol H2O.
Jensen
Chem 110 Chap 3 & 4
Page: 2
Balancing Chemical Equations
When reactants and products are both given,
chemical reactions are balanced by changing
the _____________.
Four “Easy” Rules
1. Write the unbalanced molecular equation
correctly (molecules involved with correct
molecular formulas)
C6H6 +
O2 ! CO2 +
H 2O
Example: When the following equation
C4H8O2(l) + O2(g) ! CO2(g) + H2O(g)
is balanced with the smallest possible set of
integer coefficients, the coefficient of O2 is
A.
B.
C.
D.
E.
1
2
3
5
6
2. Balance the atoms of one element.
C6H6 +
O2 !
CO2 +
H 2O
3. Balance atoms of remaining elements
C6H6 +
O2 !
CO2 +
H 2O
How many moles of water will be produced if
3 moles of C4H8O2 burn completely in air?
4. Check you work! Make sure that you use the
smallest whole numbers.
Jensen
Chem 110 Chap 3 & 4
Page: 3
Jensen
Chem 110 Chap 3 & 4
Page: 4
Patterns of Reactivity
Patterns of Reactivity
1. Combination reactions (Chapter 3)
5.
Exchange reactions (Chapter 4)
(Double Displacement or Metathesis Rxn)
Elements react to form compounds
2 Mg (s) + O2 (g) ! 2 MgO (s)
Small compounds combine to form larger ones
MgO (s) + CO2 (g) ! MgCO3 (s)
Exchange reactions only occur if there is a
driving force
a. Precipitation
Pb(NO3)2(aq) + 2 KI(aq) ! PbI2(s) " + 2 KNO3(aq)
2. Decomposition reactions (Chapter 3)
2 H2O (l) ! 2 H2 (g) + O2 (g)
b.Neutralization (weak or non-electrolyte)
CaCO3 (s) ! CaO (s) + CO2 (g)
NaOH(aq) + HCl(aq) ! NaCl(aq) + H2O(!)
3. Complete Combustion reactions (Chapter 3)
CH4 (g) + 2 O2 (g) ! CO2 (g) + 2 H2O(g)
c. Gas formation
2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq)
All hydrocarbons will produce ___________ when
How do you predict the phase of the products?
they undergo complete combustion reactions.
4. Single displacement reactions (Chapter 4)
Use Solubility Rules (Table 4.1)
How do you know what is happening?
Zn(s) + CuSO4(aq) ! ZnSO4(aq) + Cu(s)
Jensen
Chem 110 Chap 3 & 4
Page: 5
Use Net Ionic Equation
Jensen
Chem 110 Chap 3 & 4
Page: 6
Table 4.1: Aqueous Solubility of Ionic
Compounds (Table provided on your data sheet)
Soluble ionic compounds
Important exceptions
compounds containing NO3
–
None
C2H3O2
–
None
–
Cl
compounds of Ag+, Hg22+, and Pb2+
Br–
compounds of Ag+, Hg22+, and Pb2+
–
compounds of Ag+, Hg22+, and Pb2+
I
SO42–
compounds containing S
Use solubility rules to predict reactions based on
the reactants provided. ___________ drives the
reaction.
Mix MgCl2 and NaOH. What happens?
compounds of Sr2+, Ba2+, Hg22+, and Pb2+
Insoluble ionic compounds
Important exceptions
2–
Demo: Precipitation reactions
Mix KI and NaOH. What happens?
compounds of the alkali metal cations and
NH4+, Ca2+, Sr2+, and Ba2+
CO32–
compounds of NH4+,
the alkali metal cations
PO4
3–
compounds of NH4+,
Mix AgNO3 and NaCl. What happens?
the alkali metal cations
OH
–
compounds of the alkali metal cations and
NH4+, Ca2+, Sr2+, and Ba2+
Mix AgNO3 and KI. What happens?
Simplified Rules (in WATER)
1. Almost all ammonium and alkali metal salts are
soluble.
2. Most nitrates, acetates, chlorides, bromides, iodides,
and sulfates are soluble.
3. Most sulfides, carbonates, phosphates, and
hydroxides are insoluble.
Jensen
Chem 110 Chap 3 & 4
Page: 7
Jensen
Chem 110 Chap 3 & 4
Page: 8
Net Ionic Equation: involves only the ions or
molecules directly involved in the reaction
1st: Start with balanced molecular equation
Pb(NO3)2(aq) + 2 KI(aq) ! PbI2(s) + 2 KNO3(aq)
nd
2 : Dissociate all soluble strong electrolytes
(strong acids, strong bases, soluble ionic salts) to
get the Complete Ionic Equation.
Subscripts for ions in the chemical formula (in the
“Molecular” equation) become Coefficients for
those ions in the Complete Ionic Equation.
Practice Examples:
1. Which ions are spectator ions in the reaction
represented by the following molecular equation?
2AgNO3(aq) + CaCl2(aq) !" 2AgCl(s) + Ca(NO3)2(aq)
A. Ag+, Cl!, and Ca2+
B. Cl! and Ca2+
C. Ag+ and NO3!
D. Ca2+ and NO3!
E. Ca2+
2. Mixing solutions of K2SO4(aq) and BaCl2(aq)
produces an insoluble salt. What is the identity of
the spectator ions?
rd
3 : Identify spectator ions (ions that appear on both
sides of equation)
A. K+, SO42-, Ba2+, ClB. K+, SO42-
4th: Eliminate all spectator ions to get the Net Ionic
Equation (ions and molecules that directly involved
in the reaction)
C. K+, Cl-
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Jensen
Chem 110 Chap 3 & 4
Page: 9
D. Ba2+, ClE. Ba2+, SO42Chem 110 Chap 3 & 4
Page: 10
Examples of Acid–Base Neutralization
Neutralization (Acid–Base) reactions
acid + base ! salt + water
• Acids: donate H+(aq)
1. Strong acid – strong base neutralization
HCl (aq) + KOH (aq) ! KCl (aq) + H2O (l)
# Complete ionic equation
HCl (aq) !
• Bases: raises concentration of OH–(aq) ions
# Spectator ions
KOH (aq) !
• Salts: ionic compounds;
replaces H+ of acid with positive ion
e.g., HCl becomes KCl
# Net ionic equation
2. Weak acid – strong base neutralization
• Table 4.2—you"ve got this table memorized
by now, right??
HF (aq) + KOH (aq) ! KF (aq) + H2O (l)
# Complete ionic equation
# Spectator ions
# Net ionic equation
Jensen
Chem 110 Chap 3 & 4
Page: 11
Jensen
Chem 110 Chap 3 & 4
Page: 12
Practice Example:
Driving Force: Gas Formation
What is the net ionic equation for the
reaction between H3PO4 (aq) and KOH (aq)?
• direct production of gas (e.g., H2, CO2, H2S)
2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq)
• production of weak acid which decomposes to a
gas (e.g., H2CO3)
A. H3PO4 (aq) + KOH (aq)
! K3PO4 (aq) + H2O (l)
Example: Sodium Bicarbonate + HCl
Molecular equation:
B. H3PO4 (aq) + 3 KOH aq)
! K3PO4 (aq) + 3 H2O(l)
C. H+ (aq) + OH— (aq) ! H2O (l)
NaHCO3(aq)+HCl(aq) ! H2CO3(aq)+NaCl(aq)
H2CO3(aq) ! CO2(g)+H2O(l)
NaHCO3(aq)+HCl(aq)!CO2(g)+H2O(l)+NaCl(aq)
D. H+(aq) + KOH (aq) ! K+ (aq) + H2O(l)
Complete ionic equation
Na+(aq) + HCO3$ (aq) + H+(aq) + Cl$ (aq) !
E. H3PO4 (aq) + 3 OH— (aq)
Na+(aq) + Cl$ (aq) + CO2(g) + H2O(l)
! PO43— (aq) + 3 H2O(l)
Net ionic equation:
HCO3$ (aq) + H+(aq) ! CO2(g) + H2O(l)
Spectator ions:
Jensen
Chem 110 Chap 3 & 4
Page: 13
Jensen
Chem 110 Chap 3 & 4
Page: 14
Single Displacement Reactions
(Oxidation-reduction; aka, redox reactions)
Rules for determining Oxidation Numbers
1. The oxidation number of an atom of a pure
Redox reactions: Reactions where electrons
element is ______.
are transferred from one reactant to another
e.g. oxidation number for Cl2 is ___, for Fe is ___.
2. The oxidation number of a monatomic ion equals
_____________.
e.g. oxidation number for Cl– is ___, for Fe2+ is ___.
3. Some elements have the same oxidation number
in almost all their compounds, and can be used as
__________ to determine the oxidation numbers of
• Reduction reaction: gaining e–
X2 + 2 e – ! 2 X–
e.g. Cl2 + 2 e– ! 2 Cl–
other atoms in the compound.
• Oxidation reaction: losing e–
M ! Mn+ + n e–
e.g. Fe ! Fe2+ + 2 e–
4. The sum of oxidation numbers in a neutral
compound is ________; The sum of oxidation
• Oxidation and reduction are always linked
• Must be balanced: atoms/electrons/charge
Jensen
Chem 110 Chap 3 & 4
Page: 15
numbers in a polyatomic ion equals the __________
on the ion.
Jensen
Chem 110 Chap 3 & 4
Page: 16
Example: What is the oxidation number of
Mn (Manganese) in MnO4–?
Redox Reactions
• Oxidizing reagents: Elements or compounds
that oxidize the other reactant.
e.g.: O2, halogens, H2O2, HNO3, Cr2O7–, MnO4–
• Reducing agents: Elements or compounds
that reduce the other reactant.
e.g.: H2, C, metals
Practice example: What is the oxidation
state of S in H2SO4?
A.
B.
C.
D.
E.
+2
+4
+6
-2
-4
Jensen
Oxidation numbers always change in
redox reactions!
Example: Balance the reaction between solid
lead (II) oxide and ammonia gas to produce
nitrogen gas, liquid water, and solid lead.
Chem 110 Chap 3 & 4
Page: 17
Jensen
Chem 110 Chap 3 & 4
Page: 18
Examples of Single Displacement Reactions
Activity Series: predicts whether a certain metal
will be oxidized by an acid or a salt
1. Metal + Salt
Zn (s) + CuSO4 (aq) ! ZnSO4 (aq) + Cu (s)
Ionic equation:
Spectator Ions:
Net ionic equation:
What is the reducing agent (what is oxidized)?
What is the oxidizing agent (what is reduced)?
2. Metal + Acid
Zn (s) + 2 HCl (aq) ! ZnCl2 (aq) + H2 (g)
What is the reducing agent (oxidized)?
What is the oxidizing agent (reduced)?
Jensen
Chem 110 Chap 3 & 4
Page: 19
Jensen
Chem 110 Chap 3 & 4
Page: 20
Using the Activity Series
Problem Solving with Chemical Reactions
Remember: An element that is __________
in the activity series will be oxidized by the
ions of elements below it.
Basic skills:
• Avogadro"s number and the definition of mole
• How to calculate formula weight (molar mass)
• How to do the following conversions:
gram % mole
• Metals in the series will always be oxidized
gram % molecules
• What is meant by:
by the ions of elements _______
empirical formula
Zn(s) + AgNO3(aq) !
• Metals _______ H2 in series (e.g., Mg, Zn) will
be oxidized by an acid (e.g., HCl) to form H2
Na(s) + H2O(l) !
• Metals toward bottom are unreactive (e.g.,
Ag, Pt, Au); that is, the elemental form is most
stable
Chem 110 Chap 3 & 4
We use these along with balanced chemical
reactions to solve problems in chemistry
& Use balanced chemical equation to connect:
moles reactants % moles products
& Use balanced chemical equation and
conservation of mass to connect:
grams reactants % grams products
Example:
C2H5OH + 3 O2 ! 2 CO2 + 3 H2O
moles
1 mol
3 mol
2 mol
3 mol
1 mol of C2H5OH reacts with 3 mol of O2;
produces 2 mol of CO2 and 3 mol of H2O
Au(s) + H2O(g) !
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molecular formula
Page: 21
Jensen
Chem 110 Chap 3 & 4
Page: 22
Steps to solve stoichiometry problems
1. Write the balanced chemical equation (or
process)
2. Make a table & fill in given information
Recognize:
How much CO2 does your car produce?
The combustion of octane (C8H18) in the
presence of excess oxygen yields CO2 and
H2O. If 2.6 kg of octane is consumed, how
many kg of CO2 will it produce?
• what you know already and what you are
being asked for
• what connections will take you from the
knowns to unknowns
3. Make connections between measured properties
and the balanced equation
Mass
molar mass
moles
Volume concentration moles
(solutions)
P, V, T
(gases)
Ideal Gas Law
moles
4. Fill in table until you are able to solve the problem
5. Make sure your answer is REASONABLE
Jensen
Chem 110 Chap 3 & 4
Page: 23
Jensen
Chem 110 Chap 3 & 4
Page: 24
Limiting Reagents
& Reactant that is consumed completely;
& Determines the final amount of product;
& Must start with a balanced reaction.
When reactants mixed in unbalanced proportions,
some are left over (the ones in excess, unreacted)
Be sure to test all reactants!!!
• making a ham sandwich analogous to a chemical
reaction
Jensen
Chem 110 Chap 3 & 4
Page: 25
Example: If 36.6 g of C2H5OH reacts with
63.8 g of O2 to form CO2 and H2O, how
many grams of CO2 will be produced?
A.
B.
C.
D.
E.
26.0g
43.2g
58.5 g
70.4 g
100.4g
Jensen
Chem 110 Chap 3 & 4
Page: 26
Practice Example: The combustion reaction
Percent Yield
Theoretical yield: the yield of product that results
when the limiting reagent is completely consumed
Actual yield: the yield you actually get in the real
world
Percent yield:
! actual yield $
% yield = #
' 100
" theoretical yield &%
between 1.0 mole of C3H8 (g) and 1.0 mole of
O2 (g) goes to completion:
C3H8(g) + 5 O2(g) ! 3 CO2(g) + 4 H2O(g)
Which of these statements are true?
i. All of the C3H8 (g) is used up.
ii. 3.0 moles of CO2 (g) is formed
• Calculation is just one more step beyond a
standard stoichiometry calculation
• NOTE: if you get a % yield >100% something is
wrong, you"ve just created matter!??!!
Example continued:
If you obtained 50.0 g of CO2 from your
reaction, what is the percent yield?
Jensen
Chem 110 Chap 3 & 4
Page: 27
iii. 0.8 moles of H2O (g) is formed
A.
B.
C.
D.
E.
i only
ii only
i and ii only
iii only
i and iii only
Jensen
Chem 110 Chap 3 & 4
Page: 28
Practice Example: Lithium and nitrogen react
to produce lithium nitride as follows:
Scratch Paper:
6 Li (s) + N2 (g) ! 2 Li3N (s)
If 5.00 g of each reactant undergo a
reaction with a 80.5% yield, how many grams
of Li3N are obtained from the reaction?
A.
B.
C.
D.
E.
6.73
1.67
8.36
2.08
2.79
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g
g
g
g
g
Chem 110 Chap 3 & 4
Page: 29
Jensen
Chem 110 Chap 3 & 4
Page: 30